Category Archives: modeling instruction

Hi friends! I took a little hiatus from writing to focus on what is happening in my classroom. Now that I have some free time, I wanted to share about some changes I made last year and some changes I want to make next year. Vacation is over and my mind is already on next school year!

Some of you who are AMTA members know I made a big switch from composition books to packets last year in my general chemistry classes. Here are the reasons that lead me to making that switch:

Lost papers! I’m sure you all have that student whose organization system is one folder for all of his/her classes. Over the course of the year, that folder gets so jammed pack that it is actually just two folder halves making a paper sandwich. I have dug through many paper sandwiches for lost worksheets only to end up making extra copies.

Messy, incomplete and lost composition books. I found that students have trouble keeping neat composition books with the information I think they need. This comes back to lack of modeling expectations on my part. I also do not give any formal, PowerPoint or guided notes so students often do not know what they should write down.

The SBG thread needs to be pulled through everything. The thing that ties all of SBG together is the learning targets. I wanted a system that makes sure those learning targets are at the forefront of every activity.

Putting worksheets in a packet is easy but eliminating the composition book takes a lot more thought. I did not want to spoon feed my students everything but I also wanted them to be successful and gather the information they need in one place. I remembered seeing what Kelly O’Shea had come up with for her physics materials and that inspired what I came up with for chemistry.

I cannot post entire packets here because they include copyrighted AMTA materials but I am happy to share what I have if you have been trained in the Modeling Instruction pedagogy.

I will break down the anatomy of my packets that can be applied to any unit.

COVER PAGE: The cover page is where students write their learning targets and track their grades. This puts the learning targets at the forefront of the unit. Every time a student needs to write down a new learning target, they need to get out their packet.

LAB PAGES: Lab pages are the trickiest because you need to find a balance between guiding students without spoon-feeding them. It also helps to keep the formatting consistent so students know what to expect from a lab. I try to format it like a lab write-up: purpose, methods, data, discussion/conclusions. The only part that changes from lab to lab is the conclusions section. I like to put one or two questions at the end of each lab that sum up what I want students to take away.

WORKSHEET PAGES: Worksheet pages are the easiest. Just insert whatever worksheets you would hand out to students. Make sure to put the associated learning target(s) at the top though!

END PAGES: I end every packet with the same three pages: practicum, the model so far and additional notes. All of these blank pages are graph paper style like the lab pages. They provide students to write down practicum calculations, what we added to our model and anything else they don’t have room for. Since this is the first year I used packets, I found myself forgetting to put activities in so the additional notes pages ended up being a lifesaver!

These packets were great this past year! I only had to worry about making copies once a unit, students didn’t lose papers, learning targets actually got written down and I think students got more out of the lab activities. This coming school year my district is implementing a 1:1 Chromebook program. My summer project is figuring out how to make this system paperless! I think I will keep composition books for my honors students but my general chemistry and physical science students will be sticking with packets.

At this point in the year, the curriculum is getting more difficult and is building to what I call “the top of chemistry mountain.” I call stoichiometry the top of chemistry mountain because it pulls together the big picture of chemistry: chemical reactions, balanced equations, conservation of mass, moles and even gas laws! One of my students depicted the harrowing climb below:

Let’s recap the climb from Unit 7 before we jump in:

Molar masses on the periodic table are relative to 12 g of Carbon-12 or 1 mole of carbon

There are 6.02 x 10^23 particles in a mole

Empirical formulas represent the simplest ratio in which elements combine and can be calculated using mole ratios

Molecular formulas represent the actual number of atoms of each element that occur in the smallest unit of a molecule. This may be the same as the empirical formula.

This unit is long so you might want to pack a snack!

I start Unit 8 with an activity my students always beg me for from the first time they use Bunsen burners: making s’mores. Of course, those s’mores cost them some chemistry!

S’mores Stoichiometry

S’more stoichiometry is a fun and easy activity to introduce students to the idea of reaction ratios and even limiting reactants. A s’more can be made with the balanced equation:

Gm2 + 2Ch + Mm –> Gm2Ch2Mm

Where Gm is the diatomic element graham cracker, Ch is chocolate and Mm is marshmallow. Students go through a series of calculations converting between mass of ingredients and number of ingredients (mass of reactant to moles of reactant) and then to quantity of s’mores (moles of reactant to moles of product). Students even complete a limiting reactant problem when given a finite amount of each ingredient. The reward for all this math? Delicious, gooey, Bunsen burner s’mores.

Now that they have gotten the marshmallow roasting out of their systems, it is time to start the final ascent to the top of chemistry mountain!

BCA Tables

I love a lot of things about the Modeling Instruction curriculum, but BCA tables might be my favorite. If you are not familiar with BCA tables, check out the ChemEdX article I wrote here. BCA tables are an awesome way to help students think proportionally through stoichiometry problems instead of memorizing the mass-moles-moles-mass algorithm. I introduce BCA tables giving students moles of reactant or product. I add mass, percent yield, molarity, and gas volumes one by one as “add-ons” to the model.

Percent Yield Lab

The first “add-ons” are theoretical yield and percent yield. Students react solutions of sodium carbonate and calcium chloride (mass and mixed by students) to form calcium carbonate. Students gravity filter (I do not have aspirators in my room for vacuum filtration) the precipitate and dry it. While waiting for the product to dry, students calculate their theoretical yields. This calculation requires students to realize they need to convert their masses of reactants to moles before using a BCA table and then convert the moles of product from the BCA table to mass of product. After drying, students are able to calculate their percent yields and discuss why this is an important calculation and what their possible sources of error are.

Molarity

The next “add-on” to the BCA table is molarity. This can be saved for after limiting reactant, depending on how your schedule works out. Students learned about molarity back in Unit 7 but it never hurts to review before you jump into the stoichiometry. Again, the key to keeping this simple for students is molarity is only an add-on. Only moles can go in the BCA table so calculations with molarity should be done before or after the BCA table.

Limiting Reactant PhET

Now that students are stoichiometry pros when given excess of one reactant, it is time to “adjust to reality” as the Modeling curriculum says. This year, I introduced the concept of limiting reactants with the “Reactants, Products and Leftovers” PhET. Students started by making sandwiches with a BCA table and then moved on to real reactions. This activity helped students visualize what it looks like to have left over product. The key to using the PhET is to connect every example to the BCA table model. Before switching from sandwiches to actual reactions, I have a quick whiteboard meeting to introduce the term “limiting reactant.”

Limiting Reactant Practice

After the PhET, students work on the “Adjusting to Reality” worksheet from the Modeling Instruction curriculum. This worksheet starts by giving students reactant quantities in moles and then graduates them to mass values. The BCA table helps students easily pick out the limiting reactant and helps them see how much reactant is leftover and how much product is produced in one organized table.

I then have students work on a worksheet I call “All the Stoichiometry” because it has all types of problems with all levels of difficulty to make sure students can discern when to use the different tools they have collected.

Chemistry Feelings Circle

When I have a really challenging problem that I think would take too long for individual groups to solve, I hold a chemistry feelings circle. I arrange all of my seats in a tight circle and place a pile of whiteboards and markers in the middle.

Every student must sit in the circle and the class must solve the problem together by the end of the class period. I act like I am working on something else but really I am taking notes about their conversations. Once all students have signed off on the solution, they can elect delegates to present it to me. This year, I gave students a zombie apocalypse challenge problem involving the 2-step synthesis of putrescine. Students had to determine whether they could synthesize enough putrescine to disguise all of their classmates. Spoiler alert, there is not enough!

Ideal Gas Law

With limiting reactant under our their belts, it is time for another stoichiometry add-on, the last one. It is time for the ideal gas law. I return to gas laws through the molar volume of a gas lab. Students know how to convert mass and volume of solution to moles. What about gas volume (I may bump this back to the mole unit next year)? That question leads to the challenge of determining the volume of 1 mole of gas at STP. I usually use the traditional gas collection over water set-up but this year I was gifted a class set of LabQuest 2’s and I wanted to try them out. I used the Vernier “Molar Volume of a Gas” lab set-up instead.

I am not sold on this procedure but it got us the data we needed. With the molar volume of gas at a STP, we can derive PV=nRT and calculate R (the universal gas constant).

Grab-bag Stoichiometry

At the top of chemistry mountain, I give students a grab bag of stoichiometry problems. They may have to convert reactant or product mass, solution volume/molarity or gas volume to/from moles in addition to completing a BCA table. I give students a flow chart to fill in to help them sort out the process.

Unit 8 Practicum

Once students reach the top of chemistry mountain, it is time for a practicum. I use Flinn’s micro-mole rocket activity for the practicum but I leave it very open ended. I show students that hydrogen gas reacts with oxygen gas to form water and this creates enough energy to power the rocket (pipet bulb). From there, I set them loose to figure out what volume of each gas they need and where to mark their rocket so they can fill the gas volumes correctly. I also have students do some fun (not the word my students might use to describe them) stoichiometry calculations (see below).

Stoichiometry Coding Challenge

I usually end a unit with the practicum but I really wanted to work a computer coding challenge into this unit. Asking students to generalize the math they have been doing for weeks proves to be a very difficult but rewarding task.

For the coding challenge, I ask students to write a series of cumulative programs in Python that build to a stoichiometry calculator. First, students write a simple code that converts between mass and moles.

Then they write similar codes that convert between solution volume and moles and gas volume and moles. Students then combine those codes to create a calculator that converts any unit to moles. Once students have the front end of the stoichiometry calculator, they can add in coefficients. Finally, students build the back-end of the calculator, theoretical yield. You can read my ChemEdX blog post here.

By the end of this unit, students are about ready to jump off chemistry mountain! Luckily, the rest of the year is a downhill ski.

Let’s see what we added to the model so far…

The coefficients in a balanced equation represent the molar ratios in which elements and compounds react

The theoretical yield for a reaction can be calculated using the reaction ratios

The percent yield for a reaction is based on the quantity of product actually produced compared to the quantity of product that should theoretically be produced.

The reactant that runs out first is called the limiting reactant because it determines how much product can be produced

The pressure, volume, temperature and moles of an ideal gas can be related through the universal gas constant

We are just chugging along in chemistry this year. On to Unit 7! First let’s recap Unit 6:

Chemical reactions can be identified by a change in color, temperature or odor or the formation of a precipitate or a gas

Particles can rearrange during a chemical reaction but mass must be conserved (total number of particles does not change)

Chemical reactions occur in predictable patterns

It takes energy to break bonds and energy is released when bonds are formed

Exothermic reactions release heat when the chemical energy of the system is decreased. Endothermic reactions absorb heat when the chemical energy of the system is increased.

We have finally arrived at the mole! I know this ordering of units is a little strange but I have found that students do much better with stoichiometry if they are coming right off the mole unit.

Packing Peanut Challenge

The beginning of my mole unit is based on the concept of relative mass. I start by presenting students with this large bag of packing peanuts, a balance, and a small sample of packing peanuts and say “figure out how many packing peanuts are in here, you can’t open the bag.”

It takes students a few minutes to formulate a plan but eventually they realize they need to use the mass of their sample of packing peanuts to set up a proportion. This establishes the idea that we can count by massing. This technique is really useful when you have a large amount of something or when you need to count things that are very small (in the case of atoms, both!).

Relative Mass Activity

The packing peanuts challenge leads nicely into a more in-depth relative mass activity. I adapted this relative mass activity from the Modeling Instruction materials because I didn’t have any hardware but I did have paperclips, metal shot and pennies. In the activity, students are given vials with the same number of the aforementioned objects in each vial.

Students complete a series of calculations converting between mass and number of items. The activity ends with students calculating the relative masses of the items and comparing those relative masses to numbers on the periodic table. At this point, I make the connection that the atomic masses on the periodic table are all relative (first to hydrogen, now to carbon). Since scientists could not measure the mass of a single atom, a common sample size of particles was needed to compare masses of different elements: this is the mole. Right now, it does not matter how many particles are in a mole. All we need to know is the atomic mass on the periodic table is the mass of one mole of an element. Hence, we call this the molar mass.

I extend this discussion with a bean challenge. I give each group a vial of 50 white beans, a vial of 50 red beans, a vial with an unknown quantity of bean compounds (2 white beans and 1 red bean) and an empty vial. Students are given the challenge to determine how many bean compounds are in the mystery vial. This task requires students to find the mass of 2 white beans and 1 red bean (like finding the molar mass of a compound) and then set up a ratio to determine the number of bean compounds in the vial (like calculating the number of moles in a sample when given the mass). Students are generally able to then quickly make the connection between calculating the mass of a bean compound and calculating the molar mass of a chemical compound.

After completing these two activities, students can very easily move to practicing mass/mole conversion calculations.

Once students have relative mass down, we can figure out exactly how big a mole is.

Size of Mole

I start the discussion about the size of a mole by asking students to measure out a mole of water. This takes a little bit on thinking initially but eventually students remember from Unit 1 that the density of water is 1 g/mL so 1 mole of water would be equal to about 18 mL of water.

I then ask students “how many particles of water do you think make up that 18 mL by order of magnitude?” Students usually guess around the order of magnitude of one trillion to 1oo trillion. They are always very surprised to learn that they grossly underestimated. I follow up this discussion with some fun, size of a mole calculations to put that giant number in perspective. Did you know that a mole of basketballs would fit in a ball bag roughly the size of the Earth?

Students are then able to complete mole/particles conversion calculations and two-step conversion calculations. While students complete these calculations, I also have them working on the multi-day nail lab.

Nail Lab

I use the nail lab to introduce the concept of empirical formula. Students observe the reaction of an iron nail with copper (II) chloride, only they do not know which ion of copper was used. Students figure out how much copper was produced and how much chlorine was used, and then calculate the mole ratio and find the empirical formula. This lab takes 3 days (set-up, collect the precipitate, dry and measure the precipitate). Since each step does not take a whole class period, I do this in conjunction with mole/particle conversion calculations. I have also used the synthesis of magnesium oxide lab for determining an empirical formula which can be done in one class period (not counting the discussion).

Empirical and Molecular Formulas

After the nail lab, I jump right into calculating empirical and molecular formulas. For next year, I think I will make a more distinct transition from empirical to molecular formulas as this year my students had some trouble delineating the two. I use hydrogen peroxide and glucose as my poster child examples for the difference between empirical and molecular formulas.

To practice with empirical and molecular formulas, I have students play a round of whiteboard speed dating (see Kelly O’Shea’s blog) with a crime scene problem. The FBI has analyzed a white powder and they need to know if it is Tylenol (like the suspect claims) or cocaine. Students analyze the data and decided what to report back to the FBI.

Additionally, I have students work on “The Strange Case of Mole Airlines.” This activity was originally published in the Journal of Chemical Education and can be easily found with a quick Google search. This activity provides a wealth of practice with empirical formulas and also gives students the chance to form some conspiracy theories! Next year, I hope to set up a whole crime scene for students analyze!

Unit 7 Practicum

As will all units, I wrap up Unit 7 with a practicum. I had students calculate the formula of a hydrate. Students came up with the general lab procedure as a class (evaporate off the water and calculate the change in mass) and completed the experiment and calculations within their groups.

The practicum puts a wrap on Unit 7! Let’s sum up what we added to the model so far…

Molar masses on the periodic table are relative to 12 g of Carbon-12 or 1 mole of carbon

There are 6.02 x 10^23 particles in a mole

Empirical formulas represent the simplest ratio in which elements combine and can be calculated using mole ratios

Molecular formulas represent the actual number of atoms of each element that occur in the smallest unit of a molecule. This may be the same as the empirical formula.

That unit sets us up well for what I call the top of chemistry mountain, stoichiometry!

One of my favorite Modeling Instruction activities is the comparison of crystal structures to derive properties of ionic, molecular and atomic substances. The original instructions for this activity have you use the Mercury software from ​​​​The Cambridge Crystallographic Data Centre to visualize 3D crystal structures. The Mercury software is simple to use and makes it easy for students to make connections between properties like boiling and melting point and crystal structure.

The only problem with the Mercury software is it does not play nice with Chromebooks. If your school is anything like mine, you have a lot of Chromebooks. It makes sense, they are affordable, fast and durable. They just lack some of the computing power and operating system of a PC or Mac. Luckily, Chromebooks come with a ton of great apps meant for the classroom. One of these great apps is MolView. MolView is very similar to Mercury as it allows students to visualize crystal structures but it is not as intuitive to use. Here is a quick walkthrough to get you started:

Go to molview.org and get started!

Does your screen look like this? (maybe with a different compound)

Good! Now find your search bar, type in “sodium chloride” but don’t press “Enter”! See that little arrow next to your search box? Click it to get a drop-down menu like this:

Select “Crystallography Open Database.” That will give you some options like this:

From what I have found, you can just click the first one and it will give you what you need. Now you should have a unit cell of sodium chloride. You can hold down your mouse clicker and drag over the structure to rotate the structure like this:

Want a bigger crystal? We can do that. Click the “Model” drop-down menu and scroll to the bottom where you should see the options, “load 2x2x2 supercell” and “load 1x3x3 supercell.” Like this:

Let’s try a “2x2x2 supercell.”

Look how pretty that is! Go ahead an repeat with any other molecular, ionic or atomic substance. For some substances, like copper, you can just type in the name of the element, press “Enter” and the unit cell will pop up! If you try that and it doesn’t work, just search the crystallography database and it should be there.

If you are using the original Modeling Instruction worksheet, make sure to use the chemical name, not the common name of the compound when you search. Make sure for sugar, you search “sucrose” and for baking soda, use “sodium hydrogen carbonate.”

In metals, the positive core has a weaker attraction to the electrons so electrons can move more freely than in non-metals, allowing metals to conduct electricity.

Metals tend to lose electrons and become positively charged cations and non-metals tend to gain electrons and form negatively charged anions.

Ions are charged all over and attract ions of opposite charge from all directions. When ions of opposite charges are attracted to each other, they form ionic bonds. Ionic substances are bonded throughout and have high melting/boiling points.

When the electrons of two non-metal atoms are attracted to the other’s positive core, a covalent bond is formed. Molecular compounds are bonded within molecules but the molecules are only attracted to each other through intermolecular attractions. Molecular substances have lower melting/boiling points compared to ionic substances.

Molecules can be attracted to each other through induced dipole-dipole attractions and permanent dipole-dipole attractions.

Ionic compounds are named by writing the metal first and then dropping the ending of the non-metal and adding the suffix -ide.

Molecular compounds are named by using the prefixes -mono, -di, -tri, -tetra, etc. to denote how many atoms of each element are present in the compound. The first element only gets a prefix if there is more than 1. For the second element, you must drop the ending and add the suffix -ide.

Yes, that was a LARGE unit!

I kick off Unit 6 by blowing stuff up (because that’s what chemistry is, right?)

Chemical Reaction Demos

I think a unit on chemical reactions should start with some chemical reactions. Insert your favorite demos here. I like to use smashing thermite, the blue bottle, mossy zinc and hydrochloric acid and of course igniting hydrogen balloons from gas produced from the previous reaction.

I need to set up this demo in the future!

I have students observe the reactions and tell me how they know a chemical reaction occurred. By the end of the class we have a good list of macroscopic observations that tell us a chemical reaction has happened.

I then use the Zn and HCl reaction to introduce what is happening at the particle level. I have students draw out the particle models of the skeleton equation and they can see that it does not follow the law of conservation of mass. That is a big chemistry faux pas! The only way to fix this is to add more HCl to the reaction. This tells us that 1 zinc atom reacts with 2 hydrochloric acid molecules to form one molecule of hydrogen gas and one compound of zinc chloride. That sets us up nicely for balancing equations.

Balancing Equations

For balancing equations, it all comes down to practice. I start my students with balancing skeleton equations and then I have them move on to constructing their own skeletons from word problems. I have every student start by drawing the particle models to balance equations. Some students graduate from this quickly while others are always stuck to it. I just encourage students to do whatever works from them and I always leave individual whiteboards (sheet protectors with a white piece of paper) out on my desks during this unit for students who need them.

Never underestimate a student’s commitment to learning!

I try to break up the monotony of balancing equation worksheets with some games. Sometimes I do speed competitions (by volunteer only so students who aren’t super fast balancers don’t feel any extra pressure) or group games like board hockey.

Once students are comfortable with balancing equations, we can move on to classifying reactions.

Classifying Chemical Reactions

I start this new topic with a pretty standard chemical reactions types lab. Students complete a series of mini experiments that are representative of the different reaction types. I like to have 2 reactions for every reaction type. I give students the reaction type and skeleton equation for each reaction. Students must record their observations, balance the equations and draw the corresponding particle models for each reaction. In that aspect, there is some confirmation built into this lab but the goal is not predict products but to find patterns.

To whiteboard this lab, I have each group whiteboard a different reaction. We then talk through each reaction and find patterns in the similar reaction types. The key questions in whiteboard meeting are: “what is similar between the two reactions you saw of this type?” and “why do you think it is called insert reaction type here.“That helps us come up with a set of rules. The rules are far more meaningful to students when they come up with them themselves versus being given the patterns through notes.

After the lab, I have students classify the reactions on a worksheet that they already balanced the equations for and we whiteboard it the next day. After classifying chemical reactions, we move on to the last topic of the unit!

Energy and Chemical Reactions

This topic brings back an old favorite, the LOL chart! Before I introduce the new and improved LOLOL chart, I show students one of my favorite demos.

Don’t worry, I use a test tube and a Swedish Fish but you get the idea. This is a very exothermic reaction so it gets the conversation about heat and reactions started. I have students balance and classify the equation and then I draw an LOLOL chart on the board. This is where I introduce Echemical, which was foreshadowed in Unit 3. I ask students where they think chemical energy comes from and they can easily tell you,”from chemical bonds.” This is where you need to address the big misconception that energy is stored in bonds. It takes energy to break bonds and energy is released when bonds are formed. Collegeboard has a quick explanation of this misconception with some nice real-life examples like, “why is hydrogen such a good fuel source if it’s not storing lots of energy in its bonds?” You can also mention activation energy here and how some reactions need a bit of energy to get started but do not require a constant energy input to proceed (I like to use the example of burning magnesium ribbon).

After that discussion, I take the students observations about the gummy bear reaction and fill in the LOLOL chart accordingly.

The Swedish Fish (sugar) starts off at room temperature. After the reaction, the products are very hot. That heat had to come from somewhere and it wasn’t from the surroundings. That means it must have come from within the system; enter Echemical. After a while, the products cool down but the reaction is over so the chemical energy stays the same. That heat leaves the system so the reaction is exothermic. The LOLOL chart tells us that more energy was released forming new product bonds than what was used to break the original reactant bonds.

This is a good time to show an endothermic demo as well. I like ammonium nitrate and water because I use ammonium chloride and barium hydroxide later on for the practicum.

I have students try to whiteboard the LOLOL chart for this reaction and then we have a quick board meeting. All that is left then is some practice… and a practicum!

Unit 6 Practicum

For the Unit 6 practicum, I try to bring in as many learning targets as possible. I give students 2 reactions to observe: magnesium ribbon and hydrochloric acid and ammonium chloride and barium hydroxide. Students must give the signs that a chemical reaction has occurred, write the balanced equation, draw the particle models, classify the reaction and draw the LOLOL chart representing the observed temperature change.

That is it for Unit 6! It is small but mighty! Let’s take a look at the model so far…

Chemical reactions can be identified by a change in color, temperature or odor or the formation of a precipitate or a gas

Particles can rearrange during a chemical reaction but mass must be conserved (total number of particles does not change)

Chemical reactions occur in predictable patterns

It takes energy to break bonds and energy is released when bonds are formed

Exothermic reactions release heat when the chemical energy of the system is decreased. Endothermic reactions absorb heat when the chemical energy of the system is increased.

In this model building series, we last left off in Unit 4, a small but mighty chunk of curriculum. This unit introduced us to Dalton and his tiny particles called atoms. Here is what we learned:

All matter is made indestructible particles called atoms.

Different types of atoms are called elements.

All atoms of the same element are identical. Different elements have different properties.

Atoms combine chemically in simple, whole number ratios to make compounds.

I you remember, Dalton wasn’t sure how atoms combined together. It will be up to J.J. Thomson (and your students!) to answer that question!

NOTE: I do not follow the original modeling order where the mole comes next. I tried that my first year and it just seemed disjointed. I do not address the mole until right before stoichiometry.

Unit 5 is all about attractions. Bonds are not a stick or a hook that holds atoms together, they are electrostatic attractions. Bonds are also not something atoms “want”, because atoms are not people. I have wrestled with this unit for the last 3 years and I think I finally have something I like.

I kick off Unit 5 with an old favorite: the sticky tape lab.

Sticky Tape Lab

I know some people use this lab as a demo because it can be time consuming and sometimes the data are questionable but I think it is worth taking the time for. In the Sticky Tape lab, students observe the interactions between 2 charged pieces of tape and other materials including another set of charged tape, foil and paper.

The tricky part of this lab is getting the tape charged correctly. I give each group a roll of tape and tell the students to give it to the best direction follower in the group. They are usually pretty self-aware. I then make the class go through the process of laying the base tape, bottom tape and top tape down on the desk, peeling up the bottom and top tapes together, stroking the 2 pieces of tape and then quickly ripping them apart as a class. Then, I go around and check to make sure every group’s tape is properly charged by discretely holding a piece of foil to each piece of tape before I allow students to collect data. Once students collect their required data, I encourage them to experiment with items around their desks. I also encourage them to rub those items on someone’s head and then see how they are attracted to the tapes, foil and paper.

The discussion of this paradigm lab really helps put the model in students’ minds. It is all about explaining microscopic phenomena using macroscopic observations. Students can quickly guess that the tape somehow becomes charged, but the important part is what that means for our model. Simply moving atoms would not make a piece of tape charged. There must be a particle within the atom that has a charge! The Thomson Plum Pudding Model is born! Charged particles were transferred from one tape to another, making one positively charged and one negatively charged. This is also a good place to talk about Benjamin Franklin and his designation of positive and negative charges.

Students understand that the top tape and bottom tape are attracted to each other because opposites attract, but they have a hard time explaining how both tapes are attracted to the neutral paper and foil. I like the Balloons and Static Electricity PhET for this. Students can easily see that charged objects can displace the mobile negative charge in atoms to produce a partial charge. We then talk about how the electrons can move more easily in a metal because the positive core does not hold onto the electrons as tightly compared to a non-metal (soupy pudding vs sticky pudding). This explains why the tape was strongly attracted to the foil and only weakly attracted to the paper. This model also explains why metals conduct electricity and non-metals do not which can be easily demonstrated with a 9-volt battery/light bulb circuit. The discussion of electricity is a perfect lead in to conductivity testing.

Conductivity Testing

In the past, I have done conductivity testing of various atomic, molecular and ionic substances as a demonstration with a large, 110 V conductivity tester, but this year I decided to get crafty. I sacrificed a string of LED Christmas lights to make these mini conductivity testers.

These are just a simple circuit with speaker wire as the leads, 9-volt batteries as the power source and Popsicle sticks as the base. The only tricky part was getting a good connection between the battery and the wire. Aluminum foil and a lot of tape proved very useful for this. This tutorial was also very helpful. I set up 6 stations for students to rotate through to test various solids and solutions for conductivity. The LED lights worked great at showing different levels of conductivity by lighting dimly or brightly.

Students were able to classify their data into 4 categories: elemental solids that conduct electricity, elemental solids that do not conduct electricity, solutions that conduct electricity and solutions that do not conduct electricity. After students noticed that the solutions that conducted electricity contained a metal, we named these ionic compounds and the other type of compound, molecular. Since ionic compounds conduct electricity, they must be composed of charged particles. The obvious next questions is, which is negative and which is positive?

Micro-Electrolysis of Copper (II) Chloride

This micro-electrolysis is an activity I have added to this unit to introduce anions and cations. I instruct students on how to set up a very simple micro-electrolysis with aquarium tubing, a 9-volt battery and mechanical pencil lead (.7 mm or thicker works best).

After letting the electrolysis run for a few minutes, students see bubbles forming on the positive electrode and with some careful wafting, they can identify it as chlorine gas. When students look carefully, they see the negative electrode is turning a reddish-brown color. Students immediately call this rust. I always ask, “what is rust?” and the students reply “iron oxide?” I then ask, “is there any iron in the solution?” Students then realize that the reddish-brown substance cannot be rust and must be copper. Since the copper (metal) is attracted to the negative electrode, it must be positively charged. That means the chlorine (non-metal) must be negatively charged. This is when I introduce the terms “cation” and “anion.”

Patterns of Charge

The Modeling materials has a worksheet called “Predicting Formulas” which I have students complete after talking about anions and cations. This worksheet gives students a variety of ionic compounds and helps them find the patterns in which the ions combine. I always intro this worksheet with, “we know from the last unit that we can find the formulas for compounds using mass ratios” so students understand where these formulas come from. After completing and discussing this worksheet, students can identify the basic patterns of charge for the main group elements.

We have been zeroing in on ionic compounds for a little while, but it is time to zoom back out and look at molecular and atomic substances as well.

Structure with MolView

In the past, I have had access to 7 laptops that I could install the Mercury Software on to look at the structures of various ionic, molecular and atomic substances. I am at a new school this year so I had to find a ChromeBook alternative. Enter MolView. MolView is an awesome Mercury Software alternative. It does not have all the compounds that Mercury does and you have to do an advanced search in the Crystallography Open Database to get unit cell structures, but it gets the job done.

I had each student manipulating the structures on a ChromeBook and I also put the structures up on the SMARTboard so we could discuss them as a class.

I made a big deal this year about ionic compounds being bonded throughout because ions are charged spheres, meaning they attract particles of the opposite charge in every direction.

From looking at the structures, students constructed rules for classifying ionic, molecular and atomic substances. I always like to show students the structures of graphite and diamond to get the discussion started on “why structure matters?” Maybe that’s just the geologist in me!

After this activity, I have students complete the “Why Structure Matters” worksheet from the Modeling materials to relate structure to melting and boiling points.

This year, I added something new before getting to nomenclature. After talking about melting and boiling points, the next obvious place to go seemed to be intermolecular forces. In the past, I talked about how there are attractions between molecules that are not as strong as ionic and covalent bonds and in Unit 3, energy had to be put into a system to overcome these attractions to change phase, but I never gave these attractions a name.

Intermolecular Attractions (Forces)

I don’t like the term intermolecular forces so I call it intermolecular attractions (IMAs), because it is more descriptive of what is actually going on. For IMAs, I borrowed a lab from my colleague across the hall and “model”fied (that’s a thing, right?) it. Students timed the evaporation of 6 molecular substances: pentane, hexane (switching out for butane next year), ethanol, methanol, ethyl acetate and acetone.

I thought students might be bored by this lab because it is kind of like watching paint dry but they were actually very enthused about how quickly some of the substances evaporated and how they “disappeared” before their eyes. I heard some great hypotheses as the students talked about which ones would evaporate fastest: “it must have something to do with mass” and “these ones have oxygen in them and these don’t.”

Students saw that the evaporation times broke the substances into 3 groups: molecules without oxygen, molecules with oxygen but no OH group, and molecules with an OH group. I named the attractions in the first group “induced dipole-dipole attractions” and the second and third group “permanent dipole-dipole attractions” (because of the electronegative oxygen). I explained that the molecules with the OH group have a special kind of permanent dipole-dipole attraction called hydrogen bonding. I of course had to introduce the term “dipole” and we talked about why permanent dipole-dipole attractions seemed to be stronger than induced dipole-dipole attractions. I did not use the terms dispersion forces or Van der Waals forces because they are not descriptive of what is actually happening.

Thoughts before moving on to nomenclature

I think this unit is the toughest modeling unit to teach because you have to teach bonding without the Bohr model. The great thing about that is you are not breeding the “atoms want 8 electrons” misconception. Atoms don’t “want” anything, they are atoms. Bonding is all about electrostatic attractions, not a set of rules. My advice is hit this hard!

Ions are formed by gaining or losing electrons. When an ion forms it is charged all over so it attracts particles of opposite charge in all directions. This is why ionic compounds do not exist in discrete, formula units. This is also why 1 sodium atom bonds with 6 chlorine atoms but only has a +1 charge.

By contrast, molecular compounds bond within molecules because the electrons of each atom are attracted to the positive core of the other atom. This is why molecular compounds do not form lattices but instead are held together by weaker intermolecular attractions.

I highly recommend reading Beyond Appearances: Students’ misconceptions about basic chemical ideas (Kind, 2004). The ideas in this paper really helped me get the big picture of this unit.

Ionic and Molecular Nomenclature

Moving on, the last thing to hit in Unit 5 is ionic and molecular nomenclature. I also took a new approach to this topic this year and had students work more independently than usual. I created a “Chemistry Ninja Warrior” system where students had to “level up” to different types of nomenclature. Different levels earned different cool stickers.

Each level had a test (worksheet) that students had to demonstrate mastery on before they moved to the next level. The goal was for every student to reach level “ninja turtle” and more advanced students could move beyond that. Students worked independently or with other students at their level on POGIL activities to learn the nomenclature rules.

I liked that my students got to work at a pace that worked for them and I got to spend more time with the students who need 1:1 attention. I used the POGIL activities this year as is but next year I think I will edit them after seeing some of the snags my students ran into.

After spending some significant time on naming, the only thing left is a practicum!

Unit 5 Practicum

This practicum is more like a “demonstration of knowledge” than a lab challenge. Each group is given a set of tables containing names of elements and a pair of dice. How the students roll the dice determines the compound they will build. They must write the formula and name of the compound they roll. For molecular compounds, students roll the dice again to get the number of each type of atom in the molecule. I also have students construct a few rules for naming and differentiating ionic and molecular compounds. This is not my favorite practicum but it does a nice job of wrapping up the unit.

Whew! I think that is a hard unit to wrap your head around! Let’s sum up the model so far…

In metals, the positive core has a weaker attraction to the electrons so electrons can move more freely than in non-metals, allowing metals to conduct electricity.

Metals tend to lose electrons and become positively charged cations and non-metals tend to gain electrons and form negatively charged anions.

Ions are charged all over and attract ions of opposite charge from all directions. When ions of opposite charges are attracted to each other, they form ionic bonds. Ionic substances are bonded throughout and have high melting/boiling points.

When the electrons of two non-metal atoms are attracted to the other’s positive core, a covalent bond is formed. Molecular compounds are bonded within molecules but the molecules are only attracted to each other through intermolecular attractions. Molecular substances have lower melting/boiling points compared to ionic substances.

Molecules can be attracted to each other through induced dipole-dipole attractions and permanent dipole-dipole attractions.

Ionic compounds are named by writing the metal first and then dropping the ending of the non-metal and adding the suffix -ide.

Molecular compounds are named by using the prefixes -mono, -di, -tri, -tetra, etc. to denote how many atoms of each element are present in the compound. The first element only gets a prefix if there is more than 1. For the second element, you must drop the ending and add the suffix -ide.

Thanks for sticking with me through that one! Stay tuned for Unit 6: Chemical Reactions!

Quick recap from the first post of this series: I start the year with some underpinnings (scientific process skills that are necessary to survive in a Modeling classroom) activities. It is there that we establish how to build a scientific model.

Continuing my series on model building, let’s talk about Democritus.

Democritus’s Atomic Theory is the foundation for all of chemistry and is incredibly relevant today. This is where the chemistry Modeling Instruction curriculum starts. Democritus made the observation that if you break a rock into tiny pieces, those pieces are still made of rock. He then inferred that if you broke that rock into tiny particles so small we can’t see them, they would still have the same rock composition. Therefore, all matter must be made of teeny tiny indestructible particles that Democritus called atoms (I don’t use the word atom until we get to Dalton to avoid confusion about compounds). The first model of the atom was born! There are some other parts to Democritus’s model like the properties of atoms are determined by the shape of the atom but I don’t address that.

I start all of chemistry with the above story about Democritus and tell my students that this is our current model of the atom because we do not have any other evidence to tell us otherwise. Then I do the exploding can demonstration because chemistry is all about blowing things up, right?

Exploding Can Demo

The exploding can demonstration helps establish the practice of drawing particle diagrams. Students are asked to draw a particle diagram before the can is lit, while the can it lit, and when the can explodes. They come up with all sorts of explanations with their particle diagrams. Sometimes they are dead on, sometimes not. The right answer is not as important as the discussion of particles.

From the exploding can demonstration you can generate some particle diagram rules. The 4 I always have them come up with are:

Particles are represented as circles, not dots

Different particles should look different

Include a key so we know which particles are which

You don’t need more than 20 particles in your diagram

I always get at least one group that tries to represent particle motion with whooshies or arrows. When I see this I ask, “why are there lines coming off your particles?”, to which students usually reply, “because it’s a gas and gas particles are always moving.” I then ask, “do you have any evidence that particles move?”, to which students usually reply, “yeah, my 9th grade science teacher told me they do!” I followed that up with, “but how do you know?” It sometimes takes a few more questions to convince the students that particle motion is not currently in our model but we may add it later if we have evidence to support it. I do not tell students how the exploding can works here because they do not have the background knowledge to fully appreciate the chemistry. Instead I bring it back on the last day of class and have the students try to explain it again with their more robust model of the atom.

The discussion of particles and particle diagrams leads us straight into the “Mass and Change Lab.”

Mass and Change Lab

The “Mass and Change Lab” is a fairly standard conservation of mass lab. I have edited the lab so it is not exactly the same as what is in the Modeling Instruction materials but it includes a variety of chemical and physical changes that gain, lose and keep the same mass (depending on how you define the system). I have students use triple-beam balances during this lab to continue to reinforce the concept of significant figures.

After every group has collected their data, I have the class compile their data on the main board in the class. Each group writes whether the experiment gained or lost mass and if so, how much? The data will not be perfect. You can usually spot which groups forgot to account for the mass of a test tube or beaker and use it as an opportunity to talk about sources of error. Once we have established the mass change for each experiment, we whiteboard a before and after particle diagram for each mini experiment.

During this whiteboard session I ask students, “how are you going to show if the mass changed or stayed the same?” This is where students make the connection that the number of particles represents the mass. If the system gained mass, it must have gained particles. If the system lost mass, it must have lost particles. Students can then answer the questions, “where did the extra particles come from?” or “where did the particles go?” These questions can lead to a discussion on “what is a system?” and “what are open and closed systems?” After we have established particle diagrams for each mini-experiment, I ask students to come up with a definition for the Law of Conservation of Mass. The class usually comes up with something like “the total number of particles stays the same in a closed system.”

Now that we have established that the number of particles represents the mass, we can move on to density.

Mass and Volume Lab

I introduce the concept of density with a set of density balls I got from Education Innovations.

The two balls have the same mass but the smaller one feels heavier than the larger one. I ask students to account for this observation by drawing particle diagrams of both balls. I do not give this explanation the name “density” yet. We simply discuss it in terms of “the mass to volume ratio.”

Next we do the density lab. I have a few sets of aluminum cylinders and PVC cylinders of various sizes that I use for this lab. Any standard density lab kit would work. I ask students to find the relationship between mass and volume for the aluminum pieces and the PVC pieces. At this point in the year the students are well versed in finding relationships so I set the students loose to collect and graph their data. They come back with completed whiteboards and a lot to discuss.

Students quickly see that their data split into two lines so they have to calculate two slopes and write two statements of relationship. On the boards pictured above, you will notice that I have my students additionally draw in the line for water so we can determine if the pieces will sink or float (steeper slope than water will sink, a shallower slope than water will float). I also have the students represent both substances with particle diagrams so they have a quantitative and qualitative representation of density. I ask many questions throughout the board meeting like, “what would be more massive, 20 mL of aluminum or 20 mL of PVC?” Or the converse, “what would take up more space, 50 g of aluminum or 50 g of PVC?” At the end of the whiteboard discussion, we establish that the slope is the mass to volume ratio which we call “density.”

I follow up this lab with some worksheets on density adapted from the Modeling Instruction materials with qualitative (particle diagram) and quantitative (graphing and proportional reasoning) density questions.

I also give students a density practicum based off of Flinn’s “Don’t Sink the Boat” activity.

Once students are comfortable with the densities of liquids and solids, we can determine the density of a gas.

Density of a Gas Lab

The “Density of a Gas Lab” is a standard collection of gas by water displacement (see Flinn’s “Scientific Laboratory Techniques Guide” for a good diagram). The gas is CO2 generated by Alka-Seltzer and water. Outlining the procedure for this lab can be a little cumbersome but my students always get great data (though there are always a few groups that need a few tries to get there).

After students have collected their mass and volume measurements of the gas they collected and calculated the density, I have them record their data on the whiteboard in the front of the room. Immediately students notice that the density of a gas is a really small number. I have students put that number in scientific notation and compare it to the densities (in scientific notation) of liquids and solids we know of. This allows us to discuss the term “order of magnitude”. I ask students “how many orders of magnitude greater is the density of water compared to the density of carbon dioxide?” Students can easily determine water is three orders of magnitude denser. What students don’t realize is that means water is 1000 times denser than carbon dioxide! That usually catches them off guard so I ask them to represent the average densities of solids, liquids and gases in 3 particle diagrams.

Students either overthink it and want their particles diagrams to be exactly quantitatively correct or they underthink it and just draw each diagram with an arbitrarily smaller number of particles. Each group presents the reasoning behind their boards and we compare each board to the actual data. After a few comparisons, students realize that to truly represent the density of a gas, they would have to draw a fraction of a particle. Since fractions of particles do not fit our model, they settle for drawing one particle in the gas particle diagram. This representation is not congruent with many textbook particle diagrams and is a big misconception among students.

We have now learned all sorts of things about how the number and arrangement of particles affects properties of matter but we still have one burning question; how tiny are these tiny particles?

Thickness of a Thin Layer Activity

I wrap up the first chemistry unit with the “Thickness of a Thin Layer” activity from the Modeling Instruction materials. In this activity, students must determine the thickness of a piece of regular foil and the thickness of a piece of heavy duty foil using what they know about the density of aluminum (calculated in density lab).

From this activity, students can determine a minimum particle size if the aluminum foil is 1 particle thick (the heavy duty foil is about 1.5 times thicker than the regular foil, so the minimum particle size is 1/3 the thickness of the heavy duty foil). I then show students a clip from “The Ring of Truth” about particle size. The examples in the clip get the minimum particle size down even smaller. You could also drop a known volume of oleic acid in a large bowl of water, calculate the area of the circle it forms and then calculate the thickness of the layer to get a smaller minimum particle size.

I wrap up the discussion by showing students the “Scale of the Universe” applet. This site does a great job of putting the size of a particle into perspective for the students (as well as the size of the universe). Make sure to show it with the sound turned up, the music is awesome!

That is the end of the first chemistry unit! To sum up the model so far…

All matter is made of tiny, indestructible, hard sphere particles

The number and size of the particles determines the mass of the substance

The number of particles in a closed system does not change

The number of particles in a certain amount of space determines the density of the substance

Particles are really small; on the the order of 10^-9 m or 1 nanometer.